pH and buffers are crucial for maintaining life's delicate chemical balance. They regulate acidity in our bodies, keeping enzymes and cells functioning properly. Without them, our biochemistry would be chaos.
Understanding pH and buffers helps us grasp how living things maintain stability. From blood to stomachs, these concepts explain how organisms adapt to changing conditions and keep their internal environments just right.
pH and Acid-Base Concepts
Understanding pH Scale and Measurement
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- pH scale measures acidity or basicity of a solution
- Ranges from 0 to 14, with 7 being neutral
- Logarithmic scale based on hydrogen ion concentration [H+]
- pH calculated using the formula: pH = -log[H+]
- Each whole number change represents a tenfold difference in H+ concentration
- pH meter uses electrodes to measure electrical potential difference in a solution
- Litmus paper changes color based on pH (red for acidic, blue for basic)
Properties and Characteristics of Acids and Bases
- Acids increase H+ concentration in solution
- Donate protons (H+) when dissolved in water
- Examples include hydrochloric acid (HCl) and acetic acid (CH3COOH)
- Bases decrease H+ concentration in solution
- Accept protons or release hydroxide ions (OH-) when dissolved in water
- Examples include sodium hydroxide (NaOH) and ammonia (NH3)
- Arrhenius definition defines acids as H+ donors and bases as OH- donors
- Brønsted-Lowry definition expands to include proton transfer between molecules
Physiological pH and Its Importance
- Normal blood pH ranges from 7.35 to 7.45
- Slight alkalinity crucial for proper enzyme function and cellular processes
- pH below 7.35 leads to acidosis, causing respiratory and metabolic problems
- pH above 7.45 results in alkalosis, affecting nerve and muscle function
- Body maintains pH through various buffer systems (respiratory and renal)
- Different organs have varying pH levels (stomach pH 1-2, small intestine pH 6-7.5)
Buffers and Buffer Systems
Buffer Composition and Function
- Buffers resist changes in pH when small amounts of acid or base are added
- Consist of a weak acid and its conjugate base or a weak base and its conjugate acid
- Work by neutralizing added H+ or OH- ions
- Common buffer pairs include acetic acid/acetate and carbonic acid/bicarbonate
- Buffer capacity refers to the amount of acid or base a buffer can neutralize
- Effective buffers have a pKa close to the desired pH of the solution
Henderson-Hasselbalch Equation and pKa
- Henderson-Hasselbalch equation calculates pH of a buffer solution
- Formula: pH = pKa + log([A-]/[HA])
- [A-] represents concentration of conjugate base
- [HA] represents concentration of weak acid
- pKa is the negative logarithm of the acid dissociation constant (Ka)
- pKa indicates the pH at which a buffer is most effective
- Lower pKa values indicate stronger acids
Bicarbonate Buffer System in Human Physiology
- Primary buffer system in human blood
- Consists of carbonic acid (H2CO3) and bicarbonate ion (HCO3-)
- Regulated by respiratory system (CO2 levels) and kidneys (HCO3- excretion)
- Equation: CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO3-
- Maintains blood pH around 7.4
- Carbonic anhydrase enzyme catalyzes the conversion of CO2 to carbonic acid
- Respiratory system adjusts CO2 levels through breathing rate
Acid-Base Balance
Titration Curves and Buffer Capacity
- Titration curves graphically represent pH changes during acid-base reactions
- X-axis shows volume of titrant added, Y-axis shows pH of solution
- Steep regions indicate rapid pH changes, flat regions show buffer action
- Equivalence point occurs when moles of acid equal moles of base added
- Buffer capacity highest at pH = pKa of the weak acid or base
- Henderson-Hasselbalch equation predicts shape of titration curve
- Titration curves used to determine pKa values and select appropriate buffers
Maintaining Acid-Base Homeostasis in the Body
- Homeostasis involves maintaining stable internal conditions, including pH
- Respiratory system regulates CO2 levels through breathing rate
- Increased breathing rate expels more CO2, raising blood pH
- Decreased breathing rate retains more CO2, lowering blood pH
- Kidneys regulate HCO3- levels through reabsorption and excretion
- Renal tubules secrete H+ ions and reabsorb HCO3- to maintain pH balance
- Bone tissue acts as a long-term buffer by releasing or absorbing calcium ions
- Protein buffers (hemoglobin, albumin) contribute to short-term pH regulation